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istry = HT ay, , strc g's's', SEMESTER ~ IV, , —— eek, , , , , , ELECTROCHEMISTRY - II, , ee + + ee Cee - - - - -, Electrochemical cells involve interconversion ‘between electrical and chemical energy., They are divided in to two classes. The first class, i.e. electrolytic cells, involve conversion of, electrical energy to chemical energy. The other class, i.e. galvanic cells, will have conversion, of chemical energy to electrical energy., , In the present chapter, galvanic cells will be studied in detail., , The following points will revise our concepts about the galvanic cells., , 1. Galvanic cells convert chemical energy into electrical energy., 2. This occurs due to the chemical reaction taking place, in the galvanic cell., , The electrical energy is generated at the cost of the free energy of the reactants. Thus,, , the. electrical work-done_by the galvanic cell must be equal to the [decrease in the free__, , / of the- reactants | ieee eee, , | 4 The ‘electrical work done, will involve passage of current, Therefore, the chemical, reaction must transfer electrons that migrate and produce current in the external, circuit., , , , | 5. The chemical reaction must involve oxidation, reduction reactions, as these types of, , | reactions can generate electrons. :, , 6. Thus, the galvanic cell will contain two half cells. Oxidation will occur at one half cell, , | and reduction at the other. These half cells are also termed as the electrodes of the, , } galvanic cell., , | 7. The reaction of oxidation or reduction occurring at the electrode will generate, , potential known as the electrode potential., , 8. The chemical reaction taking place in the galvanic cell will be sum of the two electrode, , I Teactions., , | 9%. The potential developed by the cell will be the sum of the two half cell or electrode, Potentials,, , 10. One of the potentials developed is oxidation potential and the other potential is, teduction potential., , ELECTROCHEMICAL | CONTENTIONS, , , , , Scanned with CamScanner

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‘9° Basic Principles in Physical Chemistry (S.¥.8, 5 (nstng, saad a a 1,, 90 1 contact, between the two half cells, is indicated by a single or a, C. The internal co: ., line, as the case may be., 2+(aq) /Cu(s), Zn(s) /Zn**(aq) || Cu?*(aq /, D. The polarities are assigned to the electrodes. The en at which Oxidatio,, is asalpiied negative and the other electrode positive polarity., 2+(aq)® / Cu(s), ©Zn(s) / Zn?*(aq) || Cu?*(aq), E, An arrow, directed from the electrode of negative polarity to the electrode of Posi,, ° polarity, indicates, the flow of electrons, in the external circuit., , doy, , Docc,, , ———, , ®Zn(s) / Zn2*(aq) || Cu2*(aq) / Cu(s)®, Reversible and Irreversible Cells, Reversible Cell, , A reversible cell is defined as the cell in which the cell reaction gets reversed by, applying a potential slightly greater than the emf of the cell., , Example of a reversible cell is a Daniel cell. As the cell operates, the cell reaction is,, Zn(s) + Cu®+(aq) = Zn?+(aq) + Cu(s), , As the cell is reversible, the reaction gets reversed with potential applied is greater than, the emf of the cell., , Irreversible cell, An irreversible cell is defined as the one in which the cell reaction does not ge, reversed by applying a potential slightly greater than the emf of the cell., , Zinc and copper rods immersed in a sulphuric acid solution is an example of an, irreversible cell., As the cell operates, the reaction is :, Zn(s) + 2H*(aq) = Zn?+(aq) + Ho(g), , The gas will get evolve, cell emf, the reaction is,, , Cu(s) + 2H*(aq) = Cu®*(aq) + Ho(g), The gas, now will get liberated on the zinc rod., Reversible cells will contain reversible electrodes,, Table 1 : Reversible and Irreversible Cells, , don the copper rod. On applying a potential greater than the, , , , , , Reversible Cells Irreversible Cells, 1. The cell Teaction is reversible and|/1. The cell Teaction js irrever, ib], spontaneous sible;, , , , 2. No external Power supply is necessary | 2., to drive the cell reaction. It generates its, own emf during the cell reaction,, , , , A i, an Electrolytic cell, ¥sIs to occur in, , 3. No reaction will Start unless the | 3. When the two electrodes, electrodes are connected externally. connected some other re... 2°€ not, , , , | 4. Cell_reaction will occur_only when the | 4, Cell reaction may-oteur, \_ two ___elettrede are connected two electredés are _, , eacti :, occur at one of the electrodes on may, , , , not \° if the, RE co,, externally ~~ nnected, 5. It generates electric energy during the ]5. It does not produce electric en, cell reaction spends electric energy durin 8 It, reaction. . S cell, Itisan electrolytic cell., , externally,, , 6. Itis a galvanic cell. 6., , , , Scanned with CamScanner

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proces ves ”, sT EQUATION AND ITS IMPORTANCE, The Nernst equation is basic equation in the the treatment of galvanic cells. The, change in the free “energy pf the reactants ris equal to the electrical work done by the cell., 1f, BG is the change in the free energy of thé cell reaction it must be equal to the electrical, work done. :, AG = -EcennF , Q), [1] tis ean be related to the van’t Hoff’s reaction isotherm, AG = AG°+RT In Qp, Substituting in terms of the Ecey, , , , -Ecet DF = — EennF + RT In Qp we (2), Rearranging, we get,, Ecett = Een - [RT / nF] In Qp (8), At 298K, the above equation reduces to, 0.05916) - a ., Ecett = E%en- ( 5 ) log Qp 14], , + In this case, Qp, will indicate the equilibrium constant for the cell reaction, under, conditions of operation of the cell., , The cell potential is sum of the two electrode potentials. ~, Hence, the same equation can be applied to the electrode potential as well., - 0.05916) as, Este = E°ete~ (Gz) log Qp . 18), , Where, Qp, for the equilibrium constant for the electrode reaction, , , , Any electréde reaction will be either oxidation ot reduction. The general reaction for, the electrode can be written as, Ox+ne = Red +++ (6), , Where, Ox stands for the oxidized form of the electrode material and red stands for, the reduced form of the electrode material. When an expression for Qp is written, activities, of the species should be used. However, for all practical purposes in most of the cases the, concentration terms are employed., , The equilibrium constant for the electrode reaction will be given by, , Qe = [red] / [Ox] . ‘ 1 (7), Eete = EX te — [0.05916 / n] log Qp ‘ ++(8), Eete = E%eje — [0.05916 / n] log {[red] / [Ox]} _ (9), , This equation, is the general form of the Nernst equation and can be applied to any, , ‘ype of electrode system., Types of Electrodes, _ All types of electrodes have two common features. They will have a redox couple,, fi ©xidized form and reduced form, may be in the same state or in different states. All of, , em will have a redox reaction between the oxidized and reduced form. Nernst equation, — be used for all of them. The descrption of different types of electrodes is given in the, | mentioned below :, Presentation of the electrode, Diagram, if required, Electrode reaction, by convention, reduction reaction,, elctrode potential, using Nernst equation, , Wy», examples of the given type., “'BSc.-Basic Prin, in Poy. Chem. (Sem.-IIl & IV), , , , , , , FP eM, , Scanned with CamScanner

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pig wo's’ 93, Qe = Vag? j, EMMA/A™ = Hsieh ~Some log GD, Electrode Electrode Reaction, 1, Ags) -AgCl(s) | Clr AgCl(s) + e-—* Ag(s) + Clq), 9, Pb(s) - PbSO4(s) | sot PbSO4(s) +2e-——> Pb(s) + SO} (aq), , , , , , , , p. Redox electrode, , The speciality of this electrode is, both the oxidized and the reduced form are in the, aqueous phase and usually involve two different oxidation states of the same element., , Oxidized form : M"*(aq) ;, Reduced form : M",*(aq), Presentation +: Pt / M"\*(aq), M™*(aq), , Reaction : M®,*(aq) + (mg —ny)'e = M,*+(aq) ---(18), Potential : Qe = [ammt/ am"*] ---(19), Ept/M*,MBs* = E®ete— [0.05916 /n] log Qp ---(20), , +, 0.95916 Fae (21), , 7 ;, Epes M*(ayy ,M@g*(aq) = Epe/ M™*+>M"2* (hg —nq) 108 ayo*, , , , Electrode Electrode Reaction, , , , lL Pt | Fe(2) (aq) , Fe$*) (aq) Fe(3+) (aq)+e7-——> Fe(2+) (aq), , 2 Pt | sn) (aq) , sn(4+) (aq) Sn(#) (aq)+ 2e-—> Sn(2*) (aq), , , , , , , , 8. Pt | Gel) (aq) , Ce) (aq) Celt) (ag) +e" —= Cel (aq), _ Standard Electrode Potentials, The potential developed by an electrode is dependent on the activity of the species, involved in the electrode reaction i.e. use of the Nernst equation. The electrode chosen and, assigned zero potential is standard hydrogen electrode. Just as the emf of a cell is related to, the free energy change for the cell reaction, the standard free energy change can be related, to the standard emf of the cell and the equilibrium constant for the cell reaction, , AG° = - £2.) nF, , As the cell potential is, be applied to electrode potentials as well., , Ace = he nF (83), , The standard electrode potential of an electrode is defined as the potential developed, , by the electrode under the Sa edard conditions of unit activity of the ions involved and unit, of the gases, if involved. The standard electrode potential is characteristic of the, clectrode and will be independent of the factors such as the activity of the ions involved and, , © Pressure of the gas, if present., However, a cell is formed by combination of two electrodes. The cell potential is sum, , Of the two electrode potentials. Thus, the single electrode potential cannot be determined, , Perimen; This drawback is overcome, by conventionally fixing the potential of an, setae ma and then measuring the potentials of all other electrodes, relative to the, Seetrode with the assigned value of zero potential. Thus, it is not possible to determine the, thiolate standard electrode potential of any electrode. The values measured, are relative to, , , , , , +++(22), the sum of the two electrode potentials the same equation can, , , , , , Scanned with CamScanner